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The  Passivification  of  Iron  by 
Nitric  Acid 


A  THL5I5 


SUBMITTED  TO  THE  FACULTY  OF  THE  LELAXD  STANFORD  JR. 

UNIVERSITY  IN  PARTIAL  FULFILLMENTTft5  THE 

REQUIREMENTS  FOR  THE  DEGREE  OF 

DOCTOR  OF  PHILOSOPHY 


BY 


ELTON  MARION  HOGG 


May,  1915 


The  Passivification  of  Iron  by 
Nitric  Acid 


A  THL5I5 


SUBMITTED  TO  THE  FACULTY  OF  THE  LELAND  STANFORD  JR. 

UNIVERSITY  IN  PARTIAL  FULFILLMENT  OF  THE 

REQUIREMENTS  FOR  THE  DEGREE  OF 

DOCTOR  OF  PHILOSOPHY 


BY 

ELTON  MARION  HOGG 

May,  1915 


P. 


THE  PASSIVIFICATION  OF  IRON  BY  NITRIC  ACID 


BY   EI/TON   MARION   HOGG 

Introduction 

The  passive  state -of  certain  metals  may  be  induced  by 
methods  which  resolve  themselves  into  two  general  subdivisions : 
Passivity  produced  by  strong  nitric  acid  and  other  oxidizing 
agents,  and  that  produced  by  anodic  or  cathodic  polarization. 
The  experimental  work  described  in  the  present  paper  is  con- 
fined to  a  study  of  the  passivification  of  iron  by  nitric  acid. 
This  limitation  was  imposed  upon  the  research  in  view  of  the 
fact  that  of  all  the  former  work  done,  that  with  nitric  acid 
seemed  to  be  less  understood  and  there  appeared  to  be  some 
points  of  attack  on  this  particular  phase  of  the  problem  which 
might  result  in  the  clarification  of  our  knowledge  of  the  entire 

subject. 

Historical 

In  the  historical  discussion  in  this  paper  both  general 
methods  of  passivification  will  be  briefly  discussed  and  a 
general  statement  of  the  principal  theories  will  be  made. 

The  study  of  the  passivity  of  iron  has  engaged  the  atten- 
tion of  scientific  investigators  for  nearly  one  hundred  years 
and  yet  our  knowledge  of  the  subject  is  still  very  limited.  The 
recent  activity  in  this  field,  recorded  in  the  late  symposiums 
before  the  British  Faraday  Society, l  indicates  that  our  informa- 
tion on  the  subject  has,  in  the  past,  been  greatly  circum- 
scribed. Many  new  methods  of  attack  are  described  which 
are  distinct  departures  from  those  of  the  earlier  students, 
and  many  new  and  important  facts  have  been  discovered. 
To  explain  these  facts  many  theories  are  put  forward  which, 
in  most  cases,  are  merely  modifications  of  older  ones,  so  that 


Met.  Chern.  Eng.,  u,  12,  679  (1914). 

335896 


6i8  Elton  Marion  Hogg 

for  present  purposes  an  outline  of  the  three  most  prominent 
theories  will  be  sufficient. 

Theories  of  Passivity 

1.  The  Oxide  Film    Theory   of  Faraday.1 — According  to 
this  theory  the  surface  of  the  iron  is  either  oxidized  or  the 
superficial  particles  of  the  metal  are  in  such  relation  to  the 
oxygen  of  the  electrolyte  as  to  be  equivalent  to  oxidation. 
The  formation  of  a  layer  of  oxide  is  supposed  to  cause  passivity 
by  mechanically  hindering  the  metallic  ions  from  entering  the 
solution.     To  this  theory  Hittorf  raises  the  objection  that  no 
oxide  could  be  found  which  would  display  the  necessary  prop- 
erties.    The  destruction  of  passivity  by  elevation  of  tempera- 
ture is  also  hard  to  explain  by  this  means.     Finkelstein  states 
that  if  an  oxide  is  formed  it  must  conduct  electricity  like  the 
metal  itself.     LeBlanc  raised  the  objection  that  the  reflecting 
power  of  the  surface  in  the  active  and  passive  conditions  is 
the  same;  hence,  if  a  film  is  present,  it  must  be  of  less  than 
molecular  thickness. 

2.  The  Valency  Theory  of  Kruger -Finkelstein2  and  Muller* 
— According  to  this  theory  passivity  is  due  to  the  change 
of  the  metal  to  a  nobler  modification.     This  change  of  state 
depends  somewhat  upon  temperature.     Otherwise  the  elec- 
trochemical behavior  of  the  metal  depends  on  the  relative 
oxidation  and  reduction  potentials  of  the  electrolyte. 

3.  The    Reaction    Velocity    Theory    of   LeBlanc.4 — In    its 
general  form  this  theory  states  that  passivity  is  due  to  the  slow 
rate  of  change  at  the  anode.     The  passive  metal  sends  out 
ions  into  the  solution  very  lowly;  that  is,  the  reaction 

Fe  +  •  '   =  Fe  '  ' 

proceeds  very  slowly  because  the  ionization  of  the  metal  is 
associated  with  chemical  changes,  and,  when  these  changes 
are  slow,  passivity  occurs.  Several  hypotheses  have  been 


1  Phil.  Mag.,  (3)  9,  57  (1836). 

2  Zeit.  phys.  Chem.,  39,  104  (1902). 

3  Ibid.,  48,  577  (1904). 

4  Zeit.  Elektrochem.,  6,  472  (1900);  n,  9  (1905). 


The  Passivification  of  Iron  by  Nitric  Acid  619 

advanced  regarding  the  mechanism  of  the  reaction  and  the 
following  are  of  importance : 

a.  The  Oxygen   Charge  Hypothesis   of  Fredenhagen1   and 
Muthmann    and   Frauenberger.2 — The    cause    of    passivity    is 
sought  in  the  slow  rate  of  reaction  between  the  anode  and  the 
oxygen  liberated  there,  with  the  result  that  the  anode  becomes 
charged  with  the  gas,  or  that  a  metal-oxygen  alloy  is  formed. 
Grave  objects  to  this  hypothesis  because  it  does  not  explain 
the  fact  that  iron  may  be  passive  in  alkalies  and  when  heated 
in  nitrogen. 

b.  The  Anion  Discharge  Hypothesis* — By  this  theory  the 
main  change  at  the  anode  is  not  the  formation  of  metallic 
ions,  but  the  discharge  of  anions.     The  electrode,  when  active, 
is  supposed  to  contain  hydrogen  and  the  discharged  anion 
reacts  with  and  removes  the  hydrogen.     The  slowness  of  this 
reaction  allows  oxygen  to  accumulate  on  the  metal  rendering 
it  passive. 

c.  The  Hydrogen  Activation  Hypothesis  of  Foerster4  and 
Schmidt.5 — According  to  this  hypothesis  the  normal  condition 
of  iron  is  assumed  to  be  passive  and  it  becomes  active  under 
the  influence  of  a  catalyst.     The  reaction 

Fe  +  '  '    =   Fe '  ' 

is  reversible  and  is  a  retarded  reaction  in  both  directions. 
If  this  is  true,  the  formation  of  a  film  may  be  considered  the 
consequence,  and  not  the  cause,  of  passivity.  The  cause  of 
passivity  is  taken  as  the  absence  or  the  removal  of  hydrogen 
and  the  deposition  of  oxide  follows  as  a  result  of  the  inactivity 
of  the  metal.  Foerster  claims  that  hydrogen  or  an  alloy  of 
hydrogen  and  iron  is  the  catalyst;  but  Grave  and  Schmidt 
favor  the  view  that  hydrogen  ions  are  responsible  for  the  change. 

d.  The    Retarded    Hydration    Hypothesis    of   LeBlanc.6 — 
According  to  this  view  the  active  iron  sends  out  ions  into  the 

1  Zeit.  phys.  Chem.,  43,  i  (1903);  63,  i  (1908). 

2  Sitzungber.  bayr,,  Akad.,  34,  201  (1904). 

3  Chem.  News,  108,  249,  259,  271,  283  (1913). 

4  Abhandlungen  der  Bunsen  Gesellschaft,  2  (1909). 

5  Zeit.  phys.  Chem.,  77,  513  (1911). 

6  Lehrbuch  der  Elektrochemie,  5th  Ed.,  p.  285. 


620  Elton  Marion  Hogg 

electrolyte  and  with  metals  tending  to  become  passive,  these 
ions  combine  slowly  with  the  water 

ion  +  water  =  ion  hydrate. 

The  concentration  of  free  ions  at  the  electrode  becomes  great 
and  finally  the  potential  difference  between  the  electrode  and 
electrolyte  becomes  so  great  that  the  discharge  of  anions  and 
development  of  oxygen  begins.  This  view  is  based  on  the 
observation  of  Grave  that  when  the  ion  concentration  is 
sufficiently  large,  polarization  begins  at  both  anode  and 
cathode.  LeBlanc  considers  that  the  hydration  and  dehydra- 
tion of  ions  under  certain  circumstances  may  be  a  very  slow 
process. 

To  all  of  the  above  theories  objections  may  be  raised, 
many  of  them  seemingly  valid,  and,  on  the  other  hand,  the 
observed  phenomena  seem  in  many  respects  to  fit  one  theory 
as  well  as  another. 

Definition 

Up  to  the  present  time  no  single  definition  of  passivity 
has  been  accepted  by  the  many  investigators  who  have  busied 
themselves  with  this  problem.  In  the  present  work  the  pass- 
ivity of  iron  is  taken  to  mean  that  state  of  the  metal  in  which 
it  is  not  attacked  immediately  by  dilute  nitric  acid.  Normally 
nitric  acid  of  specific  gravity  1.250  will  instantly  and  vigor- 
ously attack  iron  and  finally  dissolve  it;  but  if  the  metal  is 
first  treated  with  acid  of  densities  from  1.590  to  1.260  the 
attack  is  delayed  for  some  time,  frequently  as  long  as  seventy- 
two  hours  and  even  longer  in  some  cases.  During  this  period 
of  inactivity  the  metal  remains  bright  and  no  apparent  action 
occurs. 

Preliminary  Work 

A  number  of  simple  hand  experiments  were  performed 
to  determine,  if  possible,  whether  there  was  a  measurable 
effect  of  the  following  factors  on  the  passivity  reaction : 

1.  Varying  concentrations  of  nitric  acid. 

2.  Different  grades  of  iron. 

3.  Presence  of  iron  salts. 

4.  Contact  with  platinum  and  zinc  wire. 


The  Passimfication  of  Iron  by  Nitric  Acid 


621 


5.  Electrolytes  other  than  nitric  acid. 

6.  Varying  periods  of  time  in  passivifying  acid. 

The  work,  while  giving  many  important  facts,  yielded 
results  so  inconsistent  that  we  decided  to  pass  directly  to  a 
set  of  carefully  controlled  reaction  velocity  experiments. 

On  the  basis  of  information  derived  from  the  preliminary 
experiments,  certain  lines  of  investigation  naturally  suggested 
themselves  as  likely  to  give  fruitful  results.  After  careful 
consideration,  it  seemed  important  to  first  investigate  the 
rate  at  which  iron  is  dissolved  by  nitric  acid  in  varying  con- 
centrations and  at  various  temperatures.  It  was  hoped  that 
the  results  of  such  investigations  might  throw  some  light 
on  the  character  of  the  reaction  in  general,  and  show  whether 
the  development  of  passivity  was  gradual,  or  whether  there 
was  a  definite  state  which  might  be  looked  upon  as  a  passive 
state,  and  also  whether  this  state  was  influenced  by  tempera- 
ture. 

Velocity  of  the  Reaction  between  Iron  and  Nitric  Acid 
The  metal  used  was  "Merck's  Pure  Iron  Wire,"  in  the  form 
of  pieces  two  inches  long  (5.08  cm)  and  bent  horse-shoe  in 
shape  so  as  to  be  readily  hung  from  a  glass  hook.     The  density 
of  the  wire  was  7.850  and  each  sample  weighed  about  180  mg. 
The  acids  were  made  up  from  "Baker's  C.  P.  Nitric  Acid, 
sp.  gr.   1.42"  and  triple  distilled  water,  the  last  distillation 
being  with  potassium  permanganate  and  potassium  hydroxide. 
In  all,  sixteen  strengths  of  acid  were  prepared  with  the  follow- 
ing concentrations  : 

TABLE  i1 


Sp.  gr. 

Grams  HNO3 
in  100  cc 

Sp.gr. 

Grams  HNO3  in 

IOO  CC 

I  .025 

4-7I50 

.240 

47.4796 

1.050 

9-4395 

.250 

49-7750 

1-075 

14.1362 

.260 

52.0884 

I  .  IOO 

18.8210 

.270 

54.4449 

1  .200 

38.8320 

-275 

55.6610 

1  .210 

40.9222 

.285 

58.0563 

1  .220 

43.0416 

.300 

61.7370 

1.230 

45.8210 

.400 

91.4200 

Physico-Chemical  Tables,  John  Castell-Evans,  Vol.  2,  839-842  (1911). 


622 


Elton  Marion  Hogg 


The  apparatus,  shown  in  Fig.  i,  consisted  of  a  constant 
temperature  bath  with  its  stirring  apparatus,  constant  level 
attachment  apparatus  for  stirring  the  acid  in  the  test  tubes, 
and  thus  preventing  the  accumulation  of  gas  bubbles  on  the 
surface  of  the  iron,  and  an  attachment  for  holding  four  test 
tubes  in  the  bath.  In  the  figure,  a  indicates  a  covering  of  wool 
packing,  b  is  a  large  battery  jar,  c  is  the  stirring  apparatus  for 
the  bath,  d  is  the  constant  level  attachment,  e  is  the  ther- 
mometer, /  is  a  pulley  which  operates  the  stirring  apparatus 


Fig.  i 

for  the  acid  in  the  test  tubes,  g  is  a  glass  rod  which  is  attached 
to  the  pulley  /  and  holding  the  iron  sample  in  a  hook  at  the 
lower  end,  h  is  a  glass  tube  serving  as  a  guide  for  g  and  entering 
the  top  of  the  test  tube  through  a  perforated  cork  stopper, 
i  is  a  Jena  glass  test  tube  containing  the  acid  and  the  iron 
sample,  /  is  the  motor  which  operates  the  stirring  apparatus 
in  the  bath  and  raises  and  lowers  the  iron  sample  twenty  times 
a  minute,  and  k  shows  the  shape  and  position  of  the  iron. 
The  method  consisted  in  taking  weighed  samples  of  iron  wire 
and  allowing  twenty-five  cubic  centimeters  of  nitric  acid  of 


The  Passimfication  of  Iron  by  Nitric  Acid  623 

known  concentration  to  act  on  them  at  a  given  temperature 
for  periods  of  fifteen,  thirty,  forty-five,  and  sixty  minutes, 
respectively.  The  samples  were  removed  at  the  end  of  each 
period,  washed,  dried  and  reweighed.  The  loss  of  weight 
was  taken  as  a  measure  of  the  amount  of  reaction.  All  samples 
were  run  in  duplicate  so  that  for  each  and  every  period  of 
time,  for  every  concentration  of  acid  used,  and  for  every  given 
temperature,  two  observations  were  possible. 
Reaction  Velocity  Equations 

The  reactions  between  iron  and  nitric  acid  are  at  present 
not  well  understood.  Some  investigators  claim  that  with  con- 
centrated acid  there  is  no  action  on  the  iron;  others  state 
that  the  strong  acid  causes  the  evolution  of  nitrogen  peroxide, 
while  dilute  acid  gives  nitric  oxide.  Mellor1  is  of  the  opinion 
that  "with  dilute  nitric  acid,  hydrogen  is  not  evolved;  but 

the  acid  is  reduced  to  ammonia With  hot  nitric  acid 

ferrous  nitrate  and  nitrogen  oxides  are  formed."  Gmeliri- 
Kraut2  states  that  in  strong  nitric  acid  some  peroxide  of 
nitrogen  is  formed,  while  nitric  oxide  passes  off  from  the  action 
with  dilute  acids,  and  in  intermediate  concentrations  mixtures 
of  these  oxides  are  evolved.  It  is,  therefore,  a  difficult  matter 
to  choose  any  one  equation  to  represent  the  reaction  in  ques- 
tion. Since,  however,  stronger  acids  act  relatively  slowly, 
the  main  reaction  to  be  considered  is  that  brought  about  by 
weaker  acids,  namely  that  which  produces  nitric  oxide  instead 
of  the  peroxide. 

The  equation,  assuming  ferric  nitrate  to  be  the  iron  salt 
produced,  is 

Fe  +  4  HNO3  =  Fe(NO3)3  +  2  H2O  +  NO. 

If  ferrous  nitrate  were  formed  instead  of  the  ferric  salt,  the 
tendency  would  be  to  produce  variations  in  the  velocity  con- 
stant in  the  opposite  direction  from  those  found,  which  only 
confirms  the  more  strongly  the  conclusions  to  be  drawn  later. 

The  formation  of  nitrogen  peroxide  along  with  nitric 
oxide  would  influence  the  variations  in  the  same  sense  as  those 


1  Mellor:  "Modern  Inorganic  Chemistry,"  485  (1912). 

2  Handbuch  der  anorganischen  Chemie,  Band  i,  Abteilung  i. 


624  Elton  Marion  Hogg 

found,  thus  tending  to  somewhat  discount  the  conclusions. 
But  evidence  to  be  given  will  show  that  this  fact  may  be  ignored, 
because,  as  will  be  proven,  the  reaction  stops  going  at  any  easily 
measurable  rate  when  nitrogen  peroxide  is  present.  For 
these  reasons  the  equation  has  been  assumed  to  be  sufficiently 
correct  for  its  purpose. 

According  to  the  equation,  56  grams  of  iron  react  with 
252  grams  of  nitric  acid,  and  i  gram  of  iron  is  dissolved  by 
4x/2  grams  of  the  acid. 

The  generally  accepted  Reaction  Velocity  Equation  for 
a  heterogeneous  system,  of  which  only  one  concentration  is 
variable1  (in  this  case  the  acid),  is 

f=K.S.(A-*), 

where  doc/dd  is  the  rate  of  solution,  K  the  reaction  velocity 
constant,  S  the  surface  exposed,  and  (A  --  x)  the  concentra- 
tion of  the  acid.  For  our  purposes,  doc  measures  the  amount, 
in  grams,  of  iron  dissolved  in  the  time  do,  and,  since,  each  gram 
of  iron  dissolved  uses  up  4x/2  grams  of  nitric  acid,  the  x  in  the 
right-hand  member  of  the  equation  must  be  multiplied  by 
4*72.  The  equation  then  becomes 


On    transformation    and    integration    this    equation  becomes 

dx 

7T-      VT-T=  K.S.<#, 
(A  —  4Vi  *) 

and 

In  A  —  In  (A  —  4V2  x)  =  K  .  S  .  0. 

Then,  obtaining  the  value  for  K,  we  have 

TT   -     _1_      /*,  A 

>.S'        (A-4V2*)' 

To  obtain  an  equation  for  the  surface,  it  is  only  necessary  to 
make  use  of  the  following  equation: 


1  On  account  of  the  great  tendency  to  simplification  in  reaction  rates,  it 
is  assumed  that  the  reaction  is  proportional  to  the  concentration  of  the  acid, 
and  not  to  its  fourth  or  other  power. 


The  Passivification  of  Iron  by  Nitric  Acid  625 

where  /  is  the  length  of  the  wire,  m  its  mass,  d  its  density, 
and  TT  has  its  usual  value.     Solving  for  S,  we  have 

_     4  irlm 

~d~' 
and 

S  = 


Substituting  this  value  in  the  Reaction  Velocity  Equation, 
we  have 

i  A 

.  In 


/47r/w   '       (A  — 4VW 

6 . 


d 

in  which  A  and  x  are  measured  in  grams  of  nitric  acid  and 
not  in  gram  equivalents.  Thus  K  represents  the  number  of 
grams  of  iron  dissolved  in  one  minute  from  one  square  centi- 
meter of  surface  under  the  given  conditions  and  m  is  the 
average  mass  of  the  sample,  i.  e., 


where  w0  is  the  original  mass  and  mi  the  mass  at  time  6.     Under 
these  conditions 

V4  irlm 
~~d~ 
represents  the  average  surface. 

Experimental  Results 

The  first  series  of  observations  was  made  at  o°  C,  and  the 
value  of  the  reaction  velocity  constant  calculated  for  each 
sample  of  iron  taken.  The  results  showed  a  distinct  decrease 
of  the  value  of  the  constant  for  increasing  concentrations  of 
acid,  while  with  the  stronger  acids  there  was  also  a  decrease 
with  increasing  time.  Observations  were  also  made  at  10°  C. 
and  20°  C.  Complete  tables  giving  concentration,  time,  and 
temperature  would  be  too  large  to  present  in  full  so  we  have 
chosen  four  representative  ones,  Table  2  gives  a  full  summary 
of  the  reaction  velocity  constants  obtained  in  the  whole  series 
of  experiments.  As  it  became  evident,  after  a  a  time,  that 


626 


Elton  Marion  Hogg 


data  beyond  certain  limits  of  concentration  would  be  super- 
fluous, we  frequently  omitted  such  observations. 


TABLE  2 
Values  of  K 


T  =  o°C 
0  =  time  in  min. 

Sp.  gr. 
of  acids 

15 

30 

45 

60 

1.025 

0.00574 

0.00661 

0.00655 

0.00680 

0.00645 

0.00685 



0.00583 

1.050 

0.00431 

0.00479 

0.00442 

0.00439 

o  .  00466 



o  .  00466 

o  .  00450 

1-075 

o  .  00368 

0.00373 

o  .  00366 

0.00364 

0.004II 

o  .  00407 

0.00424 

0.00390 

I.  100 

0.00387 

0.00305 

0.00327 

o  .  0032  i 

0.00389 

0.00338 

0.00322 

0.00324 

1.200 

0.00157 

0.00134 

0.00151 

O.OOI29 

0.00152 

0.00133 

0.00149 

O.OOI29 

1.250 

0.000825 

o  .  000736 

0.000834 

0.000814 

0.000862 

0.000797 

0.000792 

o  .  000843 

1.260 

0.000254 

0.000129 

0.0000877 

0.0000554 

0.000229 

0.000123 

0.0000946 

0.0000815 

1.270 

0.0001036 

0.0000539 

0.0000414 

0.0000320 

0.0000889 

o  .  0000584 

0.0000389 

0.0000253 

1-275 

0.0000561 

0.0000372 

0.0000238 

0.0000160 

0.0000994 

0.0000280 

0.0000256 

0.0000228 

1.300 

o  .  0000463 

0.0000156 

0.0000104 

0.00000777 

0.0000285 

0.0000126 

0.0000104 

0.00000782 

1.400 

0.00000273 

0.00000137 

0.00000091 

0.00000068 

0.00000274 

0.00000137 

0.00000091 

0.00000068 

The  Passivification  of  Iron  by  Nitric  Acid 


627 


TABLE  2al 

Values  of  K 

T  =  10°  C 

0    =  time  in  min. 

Sp.  gr. 
of  acids 

15 

30 

45 

60 

I.  100 

0.00623 

0.00569 

0.00677 

o.  0061  i 

1.200 

o  .  00402 

0.00332 

o  .  00366 

0.00292 

Entire  sample  of  iron  dis- 

I .210 

0.00396 

0.00314 

solved  in  specified  time 

o  .  00402 



I  .220 

0.00401 

0.00300 

0.00367 

0.00301 

1.230 

0.002585 

O.CKX)6l5 

0.000331 

o  .  000484 

0.000814 

0.000751 

0.000338 

0.000277 

I  .240 

0.000398 

0.000193 

0.000113 

0.0001322 

0.000384 

0.000200 

0.000124 

0.0000785 

1.250 

0.000256 

O.OOOI3O 

0.0000887 

o  .  0000607 

0.000266 

0.000128 

o  .  0000867 

o  .  0000649 

I  .260 

0.000176 

0.0000885 

o  .  0000630 

o  .  0000468 

0.000171 

O.OOOIoSl 

0.0000587 

0.0000439 

1.275 

o  .  0000880 

o  .  0000439 

0.0000291 

0.0000220 

o  .  0000880 

0.0000436 

0.0000250 

0.0000219 

1.300 

0.0000437 

0.0000218 

0.0000146 

0.00001097 

0.0000438 

O.OOOOI9I 

0.0000146 

0.00000958 

I  .400 

0.00000273 

0.00000136 

0.00000089 

0.00000068 

0.00000272 

0.00000136 

0.00000091 

0.00000068 

1  With  acids  (sp.  gr.  1.025-1.075)  the  entire  sample  of  iron  dissolved  inside 
of  15  min.  and  with  acids  (sp.  gr.  1.100-1.240)  inside  of  30  min. 


628 


Elton  Marion  Hogg 


TABUS  2&1 
Values  of  K 

T  =  20°  C 

8    =  time  in  min. 


Sp.  gr. 
of  acids 

15 

30 

45 

60 

1.250 

0.000251 

O.OOOI3I 

0.0000741 

0.0000584 

0.000278 

0.000128 

O.OOOIOOO 

o  .  0000685 

I  .260 

O.OOOI7I 

0.0000888 

o  .  0000538 

0.0000429 

O.OOOI72 

o  .  0000885 

0.0000537 

0.0000401 

1-275 

0.0000936 

o  .  000044  T    °  •  00002  94 

0.0000221 

o  .  0000883 

o  .  0000442 

0.0000313 

O.OOOO22O 

1.285 

O.OOOO828 

o  .  0000443 

0.0000276 

O.OOOO22O 

0.0000828 

0.0000495 

0.0000240 

o  .  0000207 

1.300 

0.0000382 

— 

0.0000128 

o  .  00000964 



0.0000192 

0.0000164 

0.00000273 

1.400 

0.00000545 

0.00000137 

0.00000181 

0  .  00000068 

0.00000272 

0.00000136 

0.00000091 

0  .  00000068 

In  Tables  3,  4,  5  and  6  appear  the  complete  data  for  con- 
centrations of  acid  corresponding  to  densities  1.050,  1.250, 
1.260  and  1.400,  respectively,  the  measurements  being  for 
o°  C.  These  tables  are  chosen  because  they  present  typical 
forms  of  conduct,  and  some  considerable  discussion  of  them 
will  be  given,  from  which,  it  is  hoped,  the  significance  of  the 
results  in  Table  2  will  become  quite  clear. 

TABLE  3  TABLE  4 

Sp.  Gr.  =  1.050  Sp.  Gr.  =  1.250 


T  =  o°C                                          T  =  o°C 
A  =  2.360                                            A  =  12.444 

e 

X 

m 

K 

e 

X 

m 

K 

15 

0.0376 

0.1632 

0.00431 

15 

0.0395 

0.1643 

0.000825 

0  .  0403 

o.  1606 

o  .  00466 

o  .  0408 

o.  1610 

0.000862 

30 

0.0758 

0.1454 

o  .  00479 

30 

0.0668 

o.  1489 

o  .  000736 

— 

— 

— 

0.0723 

0.1488 

0.000797 

45 

0.0973 

0.1310 

o  .  00442 

45 

0.1052 

0.1305 

0.000835 

o.  1024 

0.1320 

o  .  00466 

o.  1009 

0.1325 

o  .  000792 

6O    I    O  .  1  2  1  2 

0.1223 

0.00439 

60 

o.  1294 

0.1179 

0.000814 

0.1232 

o.  1204 

0.00450 

0.1326 

o.  1164 

o  .  000843 

1  With  acids  (sp.  gr.  1.025-1.240)  the  entire  sample  of  iron  dissolved  inside 

of  15  min. 


The  Passivification  of  Iron  by  Nitric  Acid 


629 


TABLE  5 
Sp.  Gr.  =  1.260 
T  =  o°C 
A  =  13.022 


TABLE  6 
Sp.  Gr.  =  1.400 
T  =  o°C 

A  =  22.855 


e 

OC 

m 

K 

e 

x       m 

' 

K 

15 

0.0137 

0.1767 

0.000254 

15 

o.oooi  o.  1827 

0.00000273 

0.0124 

0.1767 

0.000229 

o.oooi  o.  1822 

0.00000274 

30 

0.0139 

0.1776 

0.000129 

30 

0.0003  0.1827 

0.00000137 

0.0134 

o.  1760 

0.000123 

0.0003  0.1809 

0.00000137 

45 

0.0140  .  o.  1705 

O.OOOO88 

45 

O.OOO2   O.  1834 

0.00000091 

0.0153  0.1756 

o  .  000095  1 

0.0003  0.1842  0.00000091 

60 

0.0117  0.1733 

0.0000551 

60 

0.0003  0.1837 

0.00000068 

0.0176  0.1741  0.000081     !  o.oooi  0.1813 

o  .  00000068 

From  an  inspection  of  Table  3,  it  is  seen  that  since  the 
value  of  K  remains  practically  constant  throughout,  the  loss 
of  iron  in  the  varying  concentrations  of  acid  is  in  exact  agree- 
ment with  the  mass  action  theory  of  reaction  velocity  in  hetero- 
geneous systems  on  the  basis  of  the  equation  assumed,  and  from 
this  fact  we  may  conclude  that  there  is  no  inhibition  to  re- 
action, and,  hence,  no  tendency  towards  the  development  of 
passivity  as  time  goes  on. 

In  Table  4  there  is  a  similar  agreement  in  the  time-con- 
centration relation,  but  the  value  of  the  constant  is  about 
one-fifth  that  of  the  corresponding  value  in  the  preceding  table. 
This  means  that  in  the  stronger  acid  there  is  a  marked  de- 
crease in  the  rate  of  reaction.  Notwithstanding  this,  how- 
ever, 1.250  acid  does  not  develop  a  passivity  which  increases 
with  the  time  of  action. 

Table  5  gives  data  which  show  the  first  evidence  of  the 
gradual  development  of  the  passive  state.  The  value  of  the 
constant  for  the  fifteen-minute  period  is  about  one-third  that 
of  the  corresponding  value  in  Table  4;  and  for  the  thirty-, 
forty-five  and  sixty-minute  periods  in  the  same  concentration 
of  acid  these  values  become  still  smaller.  The  attack  on  the 
iron  is  practically  limited  to  the  first  fifteen  minutes  of  re- 
action, as  is  shown  by  the  values  of  oc,  which,  with  slight  varia- 


1  Probable  error. 


630 


Elton  Marion  Hogg 


tions,  are  practically  constant  after  that  time.  Thus,  at  the 
end  of  the  thirty-minute  period,  the  inhibiting  force  reaches 
its  maximum  value,  and  the  attack  of  the  acid  is  practically 
stopped.  At  o°  C  the  passive  point  or  what  we  will  hereafter 
call  the  " passive  break,"  occurs  somewhere  between  acid 
concentrations  of  1.250  and  1.260. 

The  passive  break  is  shown  in  the  Curve  ia,  Plate  I, 
where  the  acid  concentrations,  expressed  in  grams  of  nitric 
acid  in  twenty-five  cubic  centimeters  of  solution,  are  abscissae 
and  the  values  of  K  X  io6  are  ordinates.  The  curve  is  plotted 
from  values  of  K  obtained  at  o°  C,  for  sixty-minute  periods, 


\ 


9QOO 


rooo 


5OOO 


3000 


/ooo 


O    2    <4     6     3     IQ  12  14    /G    /8  2O  22  24 

Plate  I1 

using  data  for  all  concentrations  of  acid  in  Table  i.  The 
acid  concentration,  12.444,  corresponds  to  HNO3  (1.250). 
The  curve  is  made  up  of  three  parts,  one  (BC)  showing  the 
gradual  decrease  of  the  velocity  constant  with  increasing  con- 
centrations, the  second  (BA)  showing  the  very  rapid  drop 
of  the  velocity  constant  to  a  very  small  value,  and  OA  showing 
the  very  slow  decrease  of  the  constant  in  passivifying  acids. 
ib  shows  the  same  curve  except  that  the  value  of  K  is  multi- 
plied by  io7  instead  of  io6.  In  these  curves  identical  letters 
refer  to  identical  points.  The  latter  curve  is  introduced  to 
show  that  the  drop  in  the  velocity  constant  is  a  gradual,  although 
a  very  rapid  one,  in  the  portion  BA. 

1  Abscissae  are  grams  HNOa  per  25  cc. 


The  Passivification  of  Iron  by  Nitric  Acid  631 

The  portions  of  the  curves  represented  by  AB  show  the  pas- 
sive break.  This  point  has  the  characteristic  that  acids  having 
greater  concentrations  will  induce  passivity,  passivity  always 
being  considered  as  a  somewhat  variable  quantity,  whereas 
those  acids  having  lower  concentrations  will  activify.  Even 
so,  it  is  to  be  seen  that  neither  passivity  nor  activity  are 
perfectly  definite  states.  All  passive  iron  even  in  passivifying 
acids  is  still  undergoing  some  solution.  In  acids  more  dilute 
than  those  corresponding  to  the  passive  break,  iron  is  slowly 
and  imperfectly  activified,  while  in  the  still  lower  concentra- 
tions this  is  accomplished  more  rapidly  and  completely. 

Table  5  gives  the  results  for  acid  of  density  1.400,  the 
highest  concentration  used  in  the  experiments.  The  value 
of  the  constant  has  suffered  a  very  marked  decrease,  and  there 
is,  nevertheless,  as  the  table  shows,  a  very  slow  rate  of  solu- 
tion. The  1.400  acid  is,  therefore,  a  strongly  passivifying 
acid. 

The  results  indicated  by  the  data  in  these  tables  is  con- 
firmed and  amplified  by  the  results  given  in  the  complete  sum- 
mary of  all  results,  namely  Table  2,  and  the  complete  data 
for  the  curves  ia  and  ib  in  Plate  I  were  taken  from  this  table. 

It  is  to  be  noted  that  in  some  cases,  namely  in  those  ex- 
periments carried  on  with  acids  of  high  passivifying  power, 
the  whole  measurable  amount  of  reaction  was  over  in  the  first 
fifteen  minutes.  In  such  cases,  of  course,  the  decrease  of  the 
values  of  K  with  increasing  time  is  without  particular  signif- 
icance. 

From  the  results  given  above,  we  must  conclude  that  an 
increase  of  concentration  of  nitric  acid  inhibits  the  rate  of 
solution  of  iron  very  greatly  even  at  concentrations  which 
are  not  sufficient  to  produce  visually  complete  passivity. 
In  concentrations  of  acid  just  below  the  passive  break,  that 
is,  from  1.200  to  1.25,  it  is  also  evident  that  the  degree  of 
inhibition  increases  very  materially  with  the  time  of  action. 
In  no  case  was  the  development  of  the  inhibition  complete 
at  the  end  of  one  hour,  although  in  some  cases  the  sample  of 
iron  had  completely  dissolved  at  that  time.  When  the  density 


Elton  Marion  Hogg 


of  the  acid  used  reaches  1.260,  there  occurs  a  break  in  the  re- 
action velocity.  The  values  of  K  fall  off  at  a  far  more  rapid 
rate  than  in  more  dilute  acids,  although  even  here  the  drop 
is  not  abrupt.  Rather  the  values  of  K  decrease  at  a  measur- 
able pace  to  the  very  small  values  corresponding  to  visually 
complete  passivity.  Visually  complete  passivity  is  not  per- 
fect passivity,  but  merely  a  very  slow  rate  of  reaction. 
Rate  of  Solution  of  Iron  in  Passivifying-  Acids 
In  order  to  be  positively  sure  that  there  was  solution  of 
iron  in  passivifying  acids,  and  to  obtain  some  data  on  the 
solution  rate,  the  following  experiments  were  devised :  Weighed 
samples  of  iron  and  ten  cubic  centimeters  of  nitric  acid  (1.300) 
were  placed  in  Jena  test  tubes  and  the  tops  of  the  tubes  drawn 
out  to  a  capillary  to  prevent  evaporation.  Duplicate  sets 
were  prepared  to  allow  of  observations  over  a  period  of  twelve 
weeks.  At  the  end  of  the  periods  specified  in  the  table  the 
samples  were  removed  from  the  acid,  washed,  dried,  and  re- 
weighed.  The  data  appear  in  Table  7.  The  entire  experi- 
ment was  repeated  with  nitric  acid  (1.400)  and  Table  8  gives 
the  results  so  obtained. 

TABLE  7  TABLE  8 

Sp.  Gr.  =  1.300  Sp.  Gr.  =  1.400 

A  =  6.174  A  =  9.142 


e 

in 
weeks 

X 

m 

K 

.0 
in 
wks. 

X 

m 

K 

! 

0.0043   0.1820 

0.00279 

I   0.0043 

0.1793 

0.00166 

0.0043  o.  1813 

0.00255 

o  .  0040 

o.  1790 

0.00157 

2 

0.0062  o.  1786 

0.00187 

2 

0.0077 

0.1778 

0.00158 

0.0060  o.  1808 

0.00181 

o  .  0079 

o.  1796 

o.  00161 

3 

o  .  0086 

0.1792 

0.00174 

3  0.0123 

0.1748 

0.00168 

o  .  0088 

o  .  i  804 

0.00179 

0.0115 

o.  1748 

0.00159 

5 

0.0150  0.1751 

0.00183 

5 

0.0209 

0.1719 

0.00176 

0.0143  0.1751 

0.00174 

0.0197 

o.  1726 

0.00165 

6  i  0.0187  o.  1752 

0.00191 

6 

0.0244 

0.1678 

0.00173 

0.0177  o.  1729 

0.00183 

0.0243 

o.  1712 

0.00170 

8 

0.0258 

o.  1694 

O.OO2OI 

8 

0.0356 

o.  1619 

0.00196 

0.0259 

o.  1690 

O.OO2O4 

0.0367 

0.1651 

0.00196 

10 

0.0347 

0.1668 

0.00220 

10 

0.0452 

0.1584 

0.00197 

o  .  0340 

0.1651 

0.00218 

0.0454 

o.  1601 

0.00197 

12 

0.0413 

0.1618 

0.00222 

12 

0.0544 

0.1544 

0  .  00202 

0.0427 

o.  1614 

0.00230 

0.0547 

0.1540 

o  .  00203 

The  Passivification  of  Iron  by  Nitric  Acid 


633 


From  these  tables  it  is  evident  that  there  is  a  continuous  rate 
of  solution  of  iron  even  in  passivifying  acids.  The  results 
further  show  that  while  the  absolute  amounts  of  solution 
(values  of  x)  is  noticeably  greater  in  the  stronger  than  in  the 
less  strong  acid,  nevertheless,  this  increase  is  not  sufficiently 
great  to  make  up  for  the  theoretical  effect  (as  demanded  by 
the  mass  action  law)  of  the  increased  concentration  of  the  acid 
used.  This  accounts  for  the  fact  that  while  in  the  stronger 
acid  the  values  of  x  are  larger,  the  values  of  K  are  smaller. 
This  points,  of  course,  to  an  increased  inhibiting  effect  with 
increasing  concentrations,  even  with  these  extremely  strong 

acids. 

The  Time-Temperature-Concentration  Function 

From  the  data  given  in  Table  2,  a  number  of  cross  com- 
pilations might  be  made.  In  Plate  II  are  shown  the  values 
of  K  X  io6  plotted  against  the  time  for  five  concentrations 


6O 


Plate  II 


of  acid  at  o°,  10°  and  20°  C.  Each  individual  curve  shows 
the  decrease  in  the  velocity  constant  with  increasing  time. 
With  decreasing  concentrations  of  acid  the  curves  assume  a 


634  Elton  Marion  Hogg 

steeper  slope,  showing  a  rapid  falling  off  in  the  value  of  the 
constant  in  partially  passivifying  acids.  It  will  be  seen  that, 
in  the  case  of  acids  of  higher  concentrations,  namely,  1.400, 
1.300,  and  1.275,  the  values  of  the  constant  for  the  sixty-minute 
periods  are  nearly  the  same  at  all  temperatures,  while  for  the 
shorter  periods  the  value  increases  as  we  pass  from  o°  to  10°  C, 
and  then  remains  constant  or  decreases  slightly  from  10° 
to  20°  C.  With  1.250  acid  at  o°  C,  the  value  of  the  constant 
is  so  large  as  to  be  entirely  off  the  plate.  Some  of  the  points 
for  the  duplicates  show  considerable  variation,  but  when  we 
consider  that  the  amounts  of  iron  dissolved  are  extremely 
small,  they  are  no  greater  than  is  to  be  expected.  The  facts, 
which  are  brought  out  above,  seem  to  indicate  a  much  slower 
development  of  passivity  at  o°  C  than  occurs  at  10°  and  20°  C, 
and  we  have  the  curious  phenomenon,  between  o°  and  10°  C, 
of  a  reaction  with  a  negative  temperature  coefficient.  Thus, 
1.250  acid  shows  no  development  of  passsivity  at  o°  C.  At 
10°  C  the  constants  are  smaller  from  the  start,  and  rapidly 
decrease  with  time.  (Compare  values  for  o°  and  10°  C  in 
Plate  II.)  In  the  summary  possible  explanations  of  these 
phenomena  will  be  given.  (See  Summary  .B,  8.) 

The  Temperature  Function  of  the  Passive  Break  Concentra- 
tion 

It  has  been  previously  pointed  out  that  at  a  given  tempera- 
ture there  is  a  fairly  definite  concentration  of  acid,  above  which 
the  reaction  rate  diminishes  rapidly.  This  point  in  concen- 
tration we  have  called  the  "passive  break."  At  different 
temperatures  this  passive  break  occurs  at  different  concen- 
trations. The  following  data  show  this.  The  values  for  o°, 
10°  and  20°  C  are  taken  from  Table  2,  while  that  for  100°  C 
is  taken  from  an  experiment  to  be  described  later. 

Passive  break  density  1.260         1.230         1.250         1.300 

Temperature  o°  10°  20°  100° 

Thus  the  passive  break  concentrations  decrease  from  o°  to 
10°  C,  and  thereafter  increase. 


The  Passivification  of  Iron  by  Nitric  Acid  635 

Miscellaneous  Experiments 

In  view  of  the  fact  that  a  better  understanding  of  some 
of  the  phenomena  of  passivity  was  obtained  from  the  pre- 
ceding investigation,  it  was  considered  advisable  to  repeat 
certain  of  the  preliminary  experiments,  not  only  as  checks, 
but  also  with  the  hope  that  new  phases  of  the  problem  would 
present  themselves. 

Different  Samples  of  Iron. — Although  many  investigators 
claim  that  impurities  do  not  influence  the  passivity  reaction, 
it  was  found  to  be  nevertheless  true  that  different  grades  of 
metal  gave  different  results.  In  the  case  of  "Stubb's  Drill 
Rod,"  nitric  acid  1.300  failed  to  passivify.  This  was  also 
true  in  a  few  cases  with  "Bessemer  Steel  Rod."  ''Merck's 
Pure  Iron  Wire,"  however,  gave  better  results,  and  acid  of 
specific  gravity  1.300  produced  passivity  in  every  case.  In  all 
of  the  following  experiments  Merck's  wire  was  used  exclu- 
sively. 

Reaction  at  Higher  Temperatures. — It  was  found  that  the 
amounts  of  iron  dissolved  were  approximately  the  same  in  acid 
of  density  i  .300  at  all  temperatures  used  in  the  reaction  velocity 
experiments.  With  a  view  of  obtaining  data  at  a  still  higher 
temperature,  the  following  experiment  was  made  at  the  boiling 
point  of  the  acid.  Nitric  acid  (1.300)  was  heated  to  its 
boiling  point  (about  115.3°  C)  and  the  sample  of  iron  dropped 
into  the  hot  acid.  Violent  continuous  action  occurred  for 
some  time,  after  which  the  iron  became  partially  passive,  a 
slow  evolution  of  nitrogen  peroxide  persisting.  The  acid  was 
allowed  to  cool  slowly  and  the  tube  was  jarred  at  frequent 
intervals. 

TABLE  9 


Temperature  at  which 
Boi,ing  point  of  acid 


115.0°  C 

115.5° 
115.3° 

101.0°  C 
99-0° 

100.0° 

84.o°C 
80.0° 
83.0° 

636 


Elton  Marion  Hogg 


This  jarring  always  produced  temporary  action  for  a  few 
seconds.  As  the  tube  cooled,  a  point  was  always  reached  at 
which  jarring  would  not  produce  this  temporary  activity. 
Table  9  gives  the  results  of  the  experiment.  While  there  is 
a  great  probability  that  there  is  considerable  solution  of  iron 
at  temperatures  below  ioo°C,  we  may  conclude  that  the  passive 
state  is  fairly  stable  under  these  conditions,  and  that  the  passive 
break  will  occur  somewhere  about  100°  C  for  acid  of  this 
strength.  No  explanation  of  the  fact  that  jarring  produces 
temporary  activity  can  be  offered  at  this  time. 

Effect  of  Time  of  Immersion. — A  sample  of  iron  was  im- 
mersed in  ten  cubic  centimeters  of  nitric  acid  (1.400)  for  five 
seconds  and  then  quickly  transferred  to  ten  cubic  centimeters 
of  1.050  acid  until  active,  after  which  it  was  returned  to  the 
first  acid  for  five  seconds.  This  procedure  was  repeated  until 
the  iron  remained  passive  in  the  dilute  acid  for  a  period  of 
at  least  ten  minutes,  the  same  portions  of  the  acids  being  used. 
Sets  were  also  run  for  immersion  periods  of  fifteen,  thirty, 
and  sixty  seconds  in  the  1.400  acid  and  the  results  are  re- 
corded in  Table  10. 

TABLE  10 
6  =  5  sec.  6  =  15  sec. 


6' 

0' 

N 

i 

2 

3 

I 

2 

3 

I 

4 

2 

4 

8 

7 

9 

2 

10 

12 

20 

12 

17 

10 

3 

18 

44 

64 

26 

18 

19 

4 

53 

60 

86 

26 

17 

15 

5 

— 

10  min. 

— 

37 

21 

30 

6 

— 

— 

— 

93 

42 

29 

7 

— 

— 

— 

— 

63. 

70 

8 

— 

— 

— 

— 

10  min. 

9 

— 

— 

— 

— 

— 

— 

10 

— 

— 

— 

— 

— 

•  — 

ii 

— 

— 

— 

— 

— 

— 

The  Passivification  of  Iron  by  Nitric  Acid 


637 


6  =  30  sec. 


6  =  60  sec. 


0' 

0' 

I 

2 

3 

I 

2 

3 

I 

6 

12         6 

15 

H 

13 

2 

14 

18 

10 

17 

20 

19 

3 

17 

17      19 

H 

16 

29 

4 

10 

17      '5 

15 

18 

35 

5 

18 

14      18 

18 

17 

46 

6 

19 

26 

27 

23 

21 

50 

7 

24 

35 

54 

21 

22      45 

8 

44 

37.     23 

23 

20     55 

9 

10  min. 

— 

30 

30 

10    1   — 

— 

— 

53 

33 

— 

ii 

— 

— 

— 

46 

45 

10  min. 

6  is  the  time  of  immersion  in  1.40x3  acid,  0'  is  the  time  in  seconds 
required  to  destroy  passivity  in  1.050  acid,  and  N  is  the 
number  of  immersions.  The  results  are  taken  from  a  complete 
series  of  observations  on  ten  independent  samples  for  each  im- 
mersion period.  Of  these,  the  results  for  three  samples  only 
are  reported  here. 

Despite  the  slight  disagreement  in  some  instances,  the 
experiment  clearly  shows  that  the  time  of  immersion  in  passivi- 
fying  acids  has  a  marked  influence  on  the  degree  of  passivity. 
For  shorter  periods  of  immersion  there  is  a  decrease  in  the 
number  of  trials  necessary  to  retain  passivity  for  ten  minutes, 
while  the  longer  passivifying  immersions  give  at  the  outset 
a  longer  persistence  of  passivity  in  the  activifying  acid.  For 
these  phenomena,  we  have  no  explanation  at  present. 

Experiments  with  the  Same  Portions  and  also  with  Fresh 
Portions  of  Acids. — It  was  thought  to  be  interesting  to  find  if 
the  time  of  activation  after  passivification  was  varied  by  the 
number  of  immersions  in  the  same  portions  and  also  in  fresh 
portions  of  both  activifying  and  passivifying  acids.  The 
following  experiments  were  carried  out  under  laboratory  con- 
ditions. The  method  of  procedure  was  the  same  as  that  em- 
ployed in  the  preceding  experiment,  one  minute  passivifying 


638 


Elton  Marion  Hogg 


immersions  being  used.     The  work  was  divided  into  five  parts 
and  five  samples  of  iron  were  used  for  each  part. 

(a)  Sample  of  iron  passivified  in  same  portion  of  1.400 
acid  and  activified  in  same  portion  of  1.050  acid. 

(b)  Sample  of  iron  passivified  in  same  portion  of  1.400 
acid  and  activified  in  fresh  portions  of  1.050  acid. 

(c)  Sample  of  iron  passivified  in  fresh  portions  of  1.400 
acid  and  activified  in  same  portion  of  1.050  acid. 

(d)  Sample  of  iron  passivified  in  fresh  portions  of  1.400 
acid  and  activified  in  fresh  portions  of  1.050  acid. 

(e)  A  fresh  sample  of  iron  passivified  and  activified  in 
the  acids  already  used  in  part  (a) . 

The  results  are  given  in  Table  1 1 . 

TABLE  n 
9    =  one  minute 

0'  =  time  in  seconds  of  persistence  of  passivity  in  1.050  acid 
N  =  number  of  immersions 


N 

Values  of  e' 

(a) 

(b) 

(c] 

(d} 

(«) 

I 

15 

H 

8 

9 

16 

8 

15 

8 

ii 

ii 

12 

2 

17 

20 

14 

15    19 

13 

17 

15 

18 

— 



3 

14 

16 

19 

19 

16 

15 

18 

18 

10  min. 

4 

15 

18 

21 

21 

18 

17 

17 

19 

5 

18 

17    24 

22 

18 

20 

15 

18 

6 

23 

21 

25 

26 

15 

20 

19 

21 

7 

21 

22 

28 

30 

16 

27 

18 

20 

8 

23 

20 

32 

36 

23 

30 

20 

24 

9 

30 

30 

33 

37 

27 

39 

25 

26 

10 

53  33 

45 

39 

29 

42 

23 

29 

ii 

10  min. 

35 

52 

4i 

— 

29 

39 

12 

51 

50 

10  min. 

31 

40 

13 

67 

57 

40 

39 

H 

78 

64 

38 

48 

15 

79 

78 

48 

55 

16 

88 

93 

61 

73 

17 

102 

96 

70 

68 

18 

1  06 

1  02 

69 

84 

Not 

Not 

passive 

passive 

The  Passivification  of  Iron  by  Nitric  Acid  639 

These  data  show  that  the  difference  in  effect  produced 
by  using  the  same  portion  or  fresh  portions  of  the  passivifying 
acid  is  very  slight,  while  in  the  case  of  the  fresh  portions  of 
the  activifying  acid  the  iron  was  not  made  passive  for  a  period 
of  ten  minutes  even  after  eighteen  immersions,  while  with  used 
portions,  a  far  less  number  of  immersions  accomplished  this. 

The  results  shown  in  part  (e)  are  most  striking.  Such 
results  as  these  may  be  duplicated  at  will.  In  fact,  where 
only  two  duplicate  experiments  are  given  in  each  part,  five 
were  actually  performed.  The  interesting  point  is  that  activi- 
fying acids,  even  as  dilute  as  1.050,  sooner  or  later  lose  their 
activifying  power  when  repeatedly  used  for  this  purpose, 
while  fresh  acids  will  always  activify.  An  explanation  for 
this  will  be  given  later.  (See  Summary  B,  3.) 

Influence  of  Different  Strengths  of  Nitric  Acid  on  Passivity. 
-The  experiments  were  carried  out  under  laboratory  condi- 
tions. The  samples  of  iron  were  subjected  to  the  action  of 
different  concentrations  of  acids  which  were  known  to  passivify 
and  then  transferred  to  acids  which  would  activify.  The 
time  of  immersion  was  sixty  seconds  in  the  case  of  the  passivi- 
fying acids.  Table  12  gives  the  data  taken  from  five  separate 
determinations  in  each  case. 

From  the  table  it  appears  that  with  decreasing  concen- 
trations of  passivifying  acids  the  persistence  of  passivity  in 
the  same  activifying  acid  always  increases.  In  the  case  of 
passivifying  acids  of  the  densities  1.260  and  1.270,  it  would 
seem  that,  by  use,  these  acids  became  diluted  to  a  strength 
below  that  required  to  produce  passivity,  although  up  to  this 
time  the  persistence  of  passivity  was  greater  than  with  stronger 
acids.  A  possible  explanation  of  the  above  phenomenon  will 
be  offered  in  the  summary.  (See  Summary  B,  9.)  It  will 
also  be  noticed  that  the  more  dilute  the  activifying  acid, 
the  less  persistent  the  passivity. 

Effect  of  Contact  with  Platinum  and  Zinc. — A  piece  of 
platinum  wire  was  wound  around  the  iron  sample  and  the 
usual  procedure  followed.  The  results  given  in  Table  13  are 


640 


Elton  Marion  Hogg 
TABLE  12 


Activifying 
acids 

Passivifying  acids 

1.250 

1.400 

1.300 

1-275 

1.270 

1.260 

B'  =  time  in  seconds 

iomin. 

10  min. 

10  min. 

10  min. 

10  min. 

I  .200 

10  min. 

10  min. 

10  min. 

10  min. 

10  min. 

I  .  100 

8       9 
18     18 

15       15 
19       23 

36       39 

17       18 

43       44 

• 

10  min. 

10  min. 

10  min. 

1.050 

15     H 

17       20 

14     16 
15     18 
18     17 

23       21 

15       18 
26       30 
36       40 
52       59 
75       93 
109     116 

25       28 
5i       35 
100       97 

135     H7 
189     204 
10  min. 

39       36 
69       66 
160     138 
196     231 
Fails  to 
make 

56     75 
153   H5 

Fails  to 
make 
passive 

21        22 

10  min. 

passive 

23       20 

30       30 

53     33 
46     45 

10  min. 

for  sixty-second  immersion  periods  in  1.400  acid  and  6'  is  the 
time  in  seconds  in  1.050  acid. 

TABLE  13 
0'  =  time  in  seconds 

^64         157" 
128         138 
72  hours 

While  no  explanation  of  the  effect  of  platinum  on  the 
passive  state  can  be  given  at  this  time,  the  results  show  that 
the  nobler  metal  exerts  an  inhibiting  effect  on  the  activifica- 
tion  of  iron  in  dilute  nitric  acid,  a  fact  previously  noted  by 
both  Faraday  and  Schoenbein.  A  peculiar  phenomenon 
was  noticed  in  connection  with  the  effect  of  platinum.  The 


The  Passivification  of  Iron  by  Nitric  Acid 


641 


nobler  metal  was  wound  around  the  middle  of  the  iron  sample 
and  the  couple  placed  in  acid  in  which  the  iron  alone  would 
have  remained  active.  It  became  passive.  If  one  end  of 
the  iron  wire  was  now  touched  with  zinc,  that  end  only  became 
active  and  remained  so  as  long  as  the  zinc  was  in  contact. 
In  some  instances  after  removing  the  zinc,  the  end  formerly 
passive  became  active  for  a  moment  and  as  it  became  passive 
again  the  opposite  end  became  active.  The  alternations 
often  occurred  four  or  five  times  before  the  entire  wire  be- 
came passive.  No  explanation  of  this  phenomenon  is  offered 
at  this  time.  Zinc  always  activified  the  iron  immediately 
by  contact  in  dilute  acids  in  which  it  would  have  otherwise 
remained  passive  for  some  time. 

The  Influence  of  Iron  Salts. — In  order  to  ascertain  the 
influence  of  iron  salts  on  passivity,  the  following  experiment 
was  performed:  Samples  of  iron,  previously  weighed,  were 
placed  in  ten  cubic  centimeters  of  nitric  acid  (1.300),  together 
with  one  cubic  centimeter  of  ferric  nitrate  solution  containing 
0.0653  grams  of  iron  and  0.5304  gram  of  nitric  acid  in  each 
cubic  centimeter.  The  density  of  the  ferric  nitrate  solution 
used  was  1.4459.  The  tops  of  the  test-tubes  containing  the 
samples  were  drawn  out  to  capillaries.  After  a  specified 
time  had  elapsed  the  tubes  were  opened,  the  iron  removed, 
washed,  dried  and  reweighed. 

TABLE  14 
A  =  6.704 

6  x  m  K 


I                  0.0052 

0.1813                  0.00280 

0.0062 

o.  1804 

0.00338 

2                         0.0079 

0.1792 

0.00220 

0.0079 

o.  1760 

0.00222 

3                 o.ono 

0.1759 

o  .  00209 

0.0126 

0.1776 

O.OO236 

5 

0.0208 

o.  1710 

O.OO239 

0.0199 

0.1753 

0.00226 

6 

0.0260 

o.  1716 

0.00247 

0.0239 

o.  1697 

0.00231 

642 


Elton  Marion  Hogg 


The  values  of  K  were  calculated  and  the  data  appear  in  Table 
14,  where  6  is  in  weeks,  x  is  loss  in  grams,  A  is  the  acid  concen- 
tration in  grams  of  nitric  acid  in  each  ten  cubic  centimeters, 
and  K  is  the  reaction  velocity  constant,  which  measures  the 
amount  of  iron  dissolved  from  one  square  centimeter  of  sur- 
face in  one  week  under  laboratory  conditions.  At  the  time  of 
starting  the  experiment  it  was  noticed  that  there  was  a  rapid 
initial  action  for  a  few  seconds  which  accounts  for  the  seemingly 
large  amount  of  iron  dissolved  in  the  first  week.  The  table 
shows  a  slight  decrease  in  the  amount  of  iron  dissolved  as  com- 
pared with  Table  7  where  there  was  not  an  excess  of  iron. 
However,  this  is  not  so  great  as  to  be  very  significant  and  we 
may  conclude  that  the  presence  of  iron  salts  has  little  or  no 
effect  on  the  passivity  reaction. 

The  Influence  of  Nitrites.— -To  determine  the  influence  of 
nitrites  on  the  passivity  reaction,  a  solution  of  sodium  nitrite 
was  prepared  containing  fifty  grams  of  the  salt  in  five  hundred 
cubic  centimeters  of  solution.  Three-hundredths  of  a  cubic 
centimeter  of  the  solution  was  added  to  ten  cubic  centimeters 
of  nitric  acid  (1.400)  and  this  combination  was  used  as  the 
passivifying  agent.  The  persistence  of  passivity  was  de- 
termined as  usual  in  i  .050  acid.  The  results  appear  in  Table  1 5 . 

15 


N 

6' 

in  seconds 

N 

6'  in  seconds 

I 

13 

10 

6 

28 

44 

2 

18 

16 

7 

37 

47 

3 

20 

25 

8 

37 

57 

4 

21 

29 

9 

10  mm. 

5 

23 

35 

These  results  are  but  little  different  from  those  obtained 
without  the  addition  of  nitrites.  The  experiment  was  re- 
peated, adding  the  nitrite  solution  to  the  activifying  acid 
instead  of  the  1.400  acid.  The  influence  of  nitrites  on  the 
reaction  is  very  evident  when  added  to  the  activifying  acid, 
in  which  case  the  iron  remained  passive  for  more  than  seventy- 
two  hours  in  every  case. 


The  Passivification  of  Iron  by  Nitric  Acid  643 

Experiments  with  Nitrogen  Peroxide  and  Iron. — Our  at- 
tention was  called  to  the  fact  that  those  strengths  of  nitric 
acid,  which  were  capable  of  passivifying  iron,  gave  off,  on 
standing,  reddish  brown  fumes  of  the  higher  nitrogen  oxides 
while  activifying  acids  are  water-white,  and  it  seemed  possible 
that  this  gas  might  be  a  factor  in  the  passivity  reaction.  To 
obtain  some  data  on  this  point  and  also  to  find  whether  a 
measurable  amount  of  this  gas  was  occluded,  the  following 
experiment  was  carried  out :  Nitrogen  peroxide  was  generated 
by  the  action  of  concentrated  nitric  acid  on  copper,  the  gas 
dried  over  phosphoric  anhydride,  and  liquefied  at  atmospheric 
pressure  in  a  bulb  cooled  with  a  freezing  mixture.  The  liquid 
so  obtained  furnished  the  gas  used  in  the  experiment. 
It  was  soon  noticed  that  when  much  of  the  gas  escaped  into 
the  room,  files  and  other  iron  instruments  became  passive. 
Iron  samples  passivified  in  the  dry  gas  remained  for  many 
hours  in  dilute  nitric  acid  without  any  inclination  toward 
activification.  Since  the  metallic  surface  remained  bright 
and  also  since  exposure  to  the  air  had  little  effect,  we  were  led 
to  suspect  that  there  was  a  possibility  that  the  metal  occluded 
much  of  the  gas  and  formed  something  of  the  nature  of  a  ni- 
trogen peroxide-iron  alloy,  or  in  some  other  way  entered  into 
reaction  with  the  iron.  Following  this  idea  an  apparatus  of 
the  design  shown  in  Fig.  2  was  prepared.  The  bulb  b  was 


Fig.  2 

filled  with  "Merck's  Pure  Iron  Wire  for  Standardization" 
and  the  entire  apparatus  thoroughly  dried  by  desiccation  over 
phosphoric  anhydride  and  aeration  with  pure,  dry  air.  The 
apparatus  was  then  connected  with  a  supply  of  nitrogen  per- 
oxide and  a  stream  of  the  gas  allowed  to  flow  through  the  tube 
for  one  hour.  The  gas  was  displaced  by  dry  air  for  the  same 
length  of  time  and  the  following  data  gathered: 


644  Elton  Marion  Hogg 

Weight  of  tube,  iron  and  nitrogen  peroxide  71 . 2900  grams 

Weight  of  tube  and  iron  71 . 2674  grams 

Weight  of  nitrogen  peroxide  occluded  0.0226  grams 

Weight  of  tube  and  iron  71 . 2674  grams 

Weight  of  tube  66 . 6001  grams 

^  

Weight  of  iron  4-6673  grams 

Weight  of  nitrogen  peroxide  occluded  by  one  gram 

of  iron  =  0.0226/4.6673  =  0.00484  grams 

Immediately  after  the  dry  iron  was  removed  from  the 
tube,  and  exposed  to  air  (the  atmosphere  was  moist)  small 
liquid  globules  appeared  on  its  surface,  which  proved  to  be 
strongly  acid.  It  was  undoubtedly  either  nitric  acid  or  a 
mixture  of  this  with  nitrous  acid.  This  point  will  be  further 
investigated. 

These  results  lead  us  to  suspect  that  passivity  is  due, 
not  to  nitric  acid  itself,  but  to  nitrogen '  peroxide,  or  at  least 
to  oxides  of  nitrogen  higher  than  nitric  oxide.  Whether  the 
active  agent  is  nitrogen  peroxide  or  possibly  nitrous  acid  is 
virtually  impossible  to  determine,  because  nitrous  acid  always 
breaks  down  readily  yielding  the  oxide  and  peroxide  of  nitro- 
gen, and  the  peroxide  on  dissolving  in  water  always  yields 
nitrous  acid.  Thus,  in  all  cases,  in  aqueous  solution,  nitrous 
acid,  nitrogen  peroxide  and  litric  oxide  presumably  exist 
together. 

It  is  perfectly  possible  that  both  nitrous  acid  and  nitrogen 
peroxide  may  be  passivifying  agents.  That  nitric  oxide  is 
not  will  be  shown  in  the  following  experiment. 

Experiment  with  Nitric  Oxide  and  Iron. — Nitric  oxide  was 
prepared  by  passing  the  gases  from  the  reaction  of  dilute  nitric 
acid  on  copper  through  a  freezing  mixture  and  then  through 
water,  thus  removing  the  peroxide.  The  gas  so  obtained  had 
no  effect  on  iron  as  far  as  passivification  was  concerned. 

History  of  the  Effect  of  Nitrogen  Peroxide  on  Iron 

It  is  of  interest  in  this  connection  to  note  that  there  are 
at  least  four  instances,  in  former  work  on  the  subject  of 
passivity,  where  nitrogen  peroxide  was  noted  by  investigators. 


The  Passivification  of  Iron  by  Nitric  Acid  645 

Schoenbein1  found  that  when  the  temperature  of  nitric 
acid  (1.36)  was  raised  to  70°  C,  a  gas  was  given  off  which  he 
called  the  "deutroxide  of  nitrogen,"  and  that  up  to  this  point 
the  iron  remained  inactive. 

Herschel2  noticed  that  after  nitric  acid  (1.399)  was  re- 
peatedly used  for  passivification,  it  became  unfit  for  use.  He 
ascribes  the  reason,  "because  it  was  impregnated  with  nitrous 
gas." 

Varenne3  found  that  when  passive  iron  was  made  active 
in  vacuo,  there  was  an  evolution  of  an  orange-colored  gas 
which  he  believed  to  be  the  "peroxide  of  nitrogen." 

Grave4  found -that  iron  heated  in  nitrogen  up  to  white 
heat  was  passive  and  also  that  ionized  nitrogen  was  a  good 
passivifying  agent.  These  phenomena  are  quite  referable 
to  traces  of  oxygen  which  under  these  conditions  might  yield 
some  nitrogen  peroxide. 

From  a  consideration  of  any  of  these  four  statements,  it 
seems  strange  that  the  possibility  of  passivification  in  nitrogen 
peroxide  directly  has  never  occurred  to  investigators  long  be- 
fore this  time,  since  it  appears  to  be  the  next  logical  step  in 
each  of  the  above  instances. 

Summary 

A.  The  results  of  an  extensive  set  of  velocity  determina- 
tions of  the  reaction  between  iron  and  nitric  acid  are  given. 
These  were  carried  out  with  acids  of  densities  ranging  from 
1.025  to  1.400,  by  small  intervals,  and  the  measurements 
were  duplicated  for  three  different  temperatures,  namely, 
o°,  10°  and  20°  C.  Isolated  observations  at  higher  tempera- 
tures are  also  given.  From  these  results  are  drawn  the  follow- 
ing conclusions,  the  reference  data  being  taken  from  results 
at  o°  C.  As  these  results  do  not  differ  very  materially  from 
those  at  10°  or  20°  C,  the  conclusions  are  generally  applicable 
to  all  results. 


1  Phil.  Mag.,  (3)9,  259  (1836). 

2  Ibid.,  (3)  ii,  329  (1836) 

3  Ibid.,   (5)  9,  76  (1836). 

4  Loc.  cit. 


646  Elton  Marion  Hogg 

1.  The  value  of  the  velocity  constant  for  the  solution  of 
iron  by  nitric  acid  decreases  with  increasing  concentration 
of  acid  throughout  the  whole  range  investigated. 

2.  This  decrease  in  rate  is  not  uniform,  but  shows  a  rapid, 
though  not  vertical  drop  in  the  velocity  constant  at  the  con- 
centration corresponding  to  a  density  of  1.260. 

3.  In  all  concentrations  of  acid  of  1.260  or  greater,  the 
value  of  the  velocity  constant  shows  a  decrease  with  the  time 
in  a  given  experiment.     That  is,  the  rate  of  solution  of  iron 
is  most  rapid  at  the  start  and  falls  off  as  the  reaction  continues. 
This  is  to  be  interpreted  as  a  gradual  development  of  passivity, 
increasing  as  the  reaction  proceeds. 

4.  The  concentration  1.260  is  thus  characterized  by  two 
things:  first,  it  is  the  point,  in  concentration  of  acid,  at  which 
progressive  passivification  begins;  second,  it  is  the  point  at 
which  the  rapid  drop  in  the  reaction  rate  takes  place.     For 
want  of  a  better  term,  we  have  called  this  point  the  "passive 
break." 

5.  The  passive  break  is  not  independent  of  the  tempera- 
ture, but  seems  to  fall  somewhat  from  o°  to  10°  C,  thereafter 
to  rise  to  20°  C  and  even  to  100°  C,  as  shown  by  a  few  ex- 
periments. 

6.  Passivifying    acids,     no    matter    how    concentrated, 
bring  about  a  slow  but  steady  solution  of  iron. 

B.  Passivification  by  Nitrogen  Peroxide.  —  It  is  shown  that 
a  degree  of  passivity,  far  greater  than  any  produced  by  the 
strongest  nitric  acid,  is  brought  about  by  exposing  iron  to 
dry  nitrogen  peroxide  gas.  Nitric  oxide  is  without  effect. 
This  fact  explains  many  of  the  phenomena  of  passivification 
which  we,  as  well  as  others,  have  observed. 

i.  Passivifying  acids  are  more  or  less  yellow  or  red, 
i.  e.,  contain  some  of  the  higher  oxides  of  nitrogen,1  while 
those  acids  which  activify,  in  general,  remain  water-  white. 


1  On  account  of  the  reversibility  of  the  reaction,  N^Os  =  NC>2  +  NO, 
in  aqueous  solution,  it  is  impossible  to  state  with  certainty  whether  the  nitrogen 
peroxide  or  the  nitrous  acid  is  the  passivifying  agent,  or  whether  both  perform 
this  function.  All  statements  that  are  made  here  concerning  the  effects  of 
nitrogen  peroxide  must,  thus,  be  considered  subject  to  this  reservation. 


The  Passivifiation  of  Iron  by  Nitric  Acid  647 

2.  The  great  preponderance  of  opinion  in  the  literature 
is  to  the  effect  that  acids  of  low  concentrations  yield  nitric 
oxide  alone,  while  those  of  higher  concentration  yield  some 
nitrogen  peroxide  also.     Our  results  would  make  it  seem  prob- 
able that  those  acids  which  passivity  are  such  as  yield  nitrogen 
peroxide   in   sufficient   amounts,    while   activifying   acids   are 
those  of  such  dilution  as  do  not  yield  sufficient  amounts  of 
this  substance.     We  have  not,  as  yet,  been  able  to  investi- 
gate this  point  any  further. 

3.  Activifying  acids  (even  quite  dilute  ones  such  as  1.050) 
lose  their  activifying  power  when  repeatedly   used.     This   is 
presumably  due  to  the  accumulation  of  dissolved  nitrogen 
peroxide  (or  nitrous  acid)  as  a  result  of  such  use. 

4.  Experiments,    which   we   have   performed   and   which 
are  too  simple  to  need  description,  showed  that  acids  in  which 
iron    is    ordinarily    active,    rapidly   produce    passivity   when 
small  quantities  of  nitrogen  peroxide  are  bubbled  through 
near  to  the  iron  which  is  being  attacked. 

5.  A  similar  effect  is  produced  by  the  addition  of  nitrites. 

6.  When  iron  is  passivified  in  nitrogen  peroxide,  a  con- 
siderable, easily  weighable,  amount  of  the  peroxide  is  absorbed 
by  the  iron.     On  exposure  to  moist  air,  the  absorbed  peroxide 
combines  with  the  moisture  and  is  converted  into  drops  of  a 
strongly  acid  liquid,  presumably  nitric  acid,  probably  mixed 
with  nitrous. 

7.  While  we  have,  as  yet,  been  unable  to  determine  the 
minimum  concentration  of  nitrogen  peroxide  capable  of  pro- 
ducing noticeable  passivity,  it  seems  likely  that  but  little  is 
necessary,  since  all  iron  and  steel  articles  in  a  room  in  which 
a  little  of  the  gas  is  allowed  to  escape,  assume  a  very  con- 
siderable   degree    of    passivity.     The    observation    of    Grave 
(loc.  cit.)  that  iron  which  is  exposed  to  the  silent  discharge 
in  tubes  evacuated  from  nitrogen,  is  probably  to  be  explained 
as  due  to  the  formation  of  some  nitrogen  peroxide  in  the  tube. 

8.  Schoenbein1  records  that  when  a  bar  of  iron  is  only 
partially  immersed  in   a  passivifying  acid,   the  unimmersed 

1  Phil.  Mag.,  (3)  9,  53  (1836). 


648  Elton  Marion  Hogg 

portion  also  becomes  passive.  We  have  found  that  this  is 
wholly  due  to  exposure  to  the  higher  oxides  of  nitrogen  (nitric 
anhydride,  N2O.5  passivifies  as  well  as  nitrogen  peroxide1).  If 
the  unimmersed  portions  are  protected  from  contact  with 
the  vapors  and  gases  evolved,  they  do  not  become  passive, 
as  we  have  determined  experimentally. 

9.  It  has  been  pointed  out  that  while  nitric  acid  of  density 
1.250  shows  markedly  the  phenomenon  of  the  development 
of  passivity  with  time  at  10°  and  20°  C,  it  does  not  show  this 
at  o°  C.     In  the  light  of  the  foregoing  discussion,  the  probable 
explanation  is  that  such  acid  develops  nitrogen  peroxide  at 
10°  C,  but  does  not  do  so  at  o°  C. 

10.  Another  curious  fact  brought  out  in  the  body  of  this 
paper  is  that  the  more  dilute  the  passivifying  acid  the  greater 
is  the  persistence  of  the  passivity  produced  when  it  is  brought 
into  the  activifying  acid.     No  very  positive  explanation  of 
this   fact   seems   possible   without   further   investigation.     It 
may  be,  however,  that,  owing  to  the  more  gradual  develop- 
ment of  the  passive  state  in  the  weaker  acids,  time  is  offered 
for  a  deeper  penetration  of  the  iron  by  the  gas,  or  by  whatever 
condition  constitutes  the  passive  state. 

C.  The  best  conception  of  the  passive  state,  as  induced 
by  nitric  acid,  which  we  are  able  to  formulate,  as  a  result  of 
our  investigations,  is  somewhat  as  follows: 

1.  The  passive  state  is  not  a  definite  one.     There  maybe 
an  indefinitely  great  number  of  degrees  of  passivity,  taking  the 
rate  of  solution  of  iron  as  a  measure  of  the  degree  of  passivity. 

2.  The  passive  state  seems  to  be  the  result  of  an  equi- 
librium between  iron  and  nitrogen  peroxide.     Iron  is  capable 
of  absorbing  nitrogen  peroxide  from  any  solution  in  which 
it  is  being  produced,  and  the  rate  of  reaction  is  thereby  in- 
hibited.    The    degree    of    inhibition    will,    therefore,    be    de- 
termined by  the  concentration  of  the  nitrogen  peroxide  which 
the  reaction  itself  is  capable  of  maintaining.     In  strong  acid 
this  concentration  is  great,   and  in  weak  acids  it  is  small, 
or  perhaps  zero.     Thus,  when  a  piece  of  iron,  which  has  ad- 
justed its  rate  of  solution  to  a  strong  passivifying  acid,  is 

1  R.  Weber:  Jour,  prakt.  Chem.,  (2)  6,  342  (1873). 


The  Passivification  of  Iron  by  Nitric  Acid  649 

brought  into  a  more  dilute  one,  it  is  charged  with  a  higher 
concentration  of  nitrogen  peroxide  than  can  be  maintained 
in  the  more  dilute  acid,  and  a  new  adjustment  takes  place, 
whereby  the  iron  slowly  gives  off  some  of  its  store  of  inhibiting 
peroxide.  This  results  in  an  increased  rate  of  solution,  which 
is  what  is  usually  called  "activification."  The  slow  develop- 
ment of  activity  in  such  cases  is  thus  explained,  as  is  also  the 
fact  that  the  "lag"  in  activification  is  less  the  more  dilute 
the  activifying  acid. 

3.  Since  iron  passivified  by  dry  nitrogen  peroxide  gas  is 
much  more  persistently  passive  in  activifying  acids  than  that 
which   is    passivified    in  nitric  acid,  it  would  seem  probable 
that  the  amount  of  nitrogen  peroxide  adsorbed  by  iron,  from 
even  very  concentrated  nitric  acid,  is  relatively  quite  small. 

4.  What   has    been    called    the    "passive  break"  in  the 
concentration  of  nitric  acid  is,  from  our  point  of  view,  to  be 
looked  upon  as  that  concentration  at  which  the  reaction  be- 
gins to  develop  relatively  large  amounts  of  nitrogen  peroxide. 
At    concentrations    below    the  passive  break,  it  would  seem 
probable    that    nitrogen   peroxide    in  decreasing    amounts  is 
still  being  developed,  since  the  values  for  the  velocity  constants 
continue    to    increase    with    dilution,  although,  as    has  been 
pointed  out,  it  is  difficult  to  say  how  much  of  this  effect  is 
due  to   increased  electrolytic  dissociation   of   the  acid.     We 
have  not,  as  yet,  been  able  to  investigate  this  point  any  fur- 
ther, but  hope  to  do  so  at  some  future  time. 

5.  As  to  why  nitrogen  peroxide  should  produce  passivity, 
or  as  to  what  the  ultimate  mechanism  of  the  process  is,  we 
do  not  feel  that  we  can  say  much.     Whether  passive  iron  is 
a  solid  solution  of  nitrogen  peroxide  in  iron,  or  whether  the 
result  of  the  action  of  the  peroxide  on  the  ions  is  the  reversible 
production    of    some    highly    oxidized  condition  of  iron,  are 
questions  which  we  do  not  feel  competent  to  discuss.      We 
may  state,   however,   that  we  have  observed  nothing  which 
seems  to  indicate  the  existence  of  anything  like  a  true  gas 
film  in  any  case. 

Laboratory  of  Physical  Chemistry 
Stanford  University 
May, 


U.  C.  BERKELEY  LIBRARII 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


